What you'll learn
This revision guide covers everything you need to know about atomic structure for AQA GCSE Physics. You'll understand the composition of atoms, how our model of the atom has developed through scientific discovery, and the structure and behaviour of the nucleus. This topic forms the foundation for understanding radioactivity and is essential for answering Paper 1 questions worth 4-6 marks.
Key terms and definitions
Atom — the smallest particle of an element that retains its chemical properties, consisting of a nucleus surrounded by electrons.
Nucleus — the tiny, dense, positively charged central core of an atom containing protons and neutrons.
Proton — a positively charged subatomic particle found in the nucleus with a relative mass of 1 and relative charge of +1.
Neutron — a neutral subatomic particle found in the nucleus with a relative mass of 1 and relative charge of 0.
Electron — a negatively charged subatomic particle that orbits the nucleus in energy levels with a relative mass of 1/2000 (or approximately 0) and relative charge of -1.
Mass number (A) — the total number of protons and neutrons in an atom's nucleus.
Atomic number (Z) — the number of protons in an atom's nucleus, which defines the element.
Isotope — atoms of the same element with the same number of protons but different numbers of neutrons.
Core concepts
Structure of the atom
Atoms have a radius of approximately 1 × 10⁻¹⁰ metres (0.1 nanometres). Despite this tiny size, most of an atom is empty space. The nucleus at the centre has a radius of about 1 × 10⁻¹⁴ metres — approximately 10,000 times smaller than the atom itself.
The atom contains three types of subatomic particle:
- Protons: located in the nucleus, relative mass = 1, relative charge = +1
- Neutrons: located in the nucleus, relative mass = 1, relative charge = 0
- Electrons: orbit the nucleus in shells (energy levels), relative mass = 1/2000, relative charge = -1
Atoms are electrically neutral because they contain equal numbers of protons and electrons. The positive charges and negative charges cancel out.
Nearly all the mass of an atom is concentrated in the nucleus because protons and neutrons are approximately 2000 times heavier than electrons. The nucleus contains more than 99.9% of the atom's mass but occupies less than 1/10,000th of its volume.
Atomic number and mass number
Every element is defined by its atomic number. This is why carbon is always carbon — it always has 6 protons.
The standard notation for representing atoms shows:
- Mass number (A) at the top left
- Atomic number (Z) at the bottom left
- Element symbol
For example: ¹²₆C represents a carbon atom with mass number 12 and atomic number 6.
To find the number of neutrons in an atom:
Number of neutrons = Mass number - Atomic number
For ¹²₆C: neutrons = 12 - 6 = 6
In a neutral atom, the number of electrons equals the number of protons (which equals the atomic number). Therefore carbon has 6 electrons.
Isotopes
Isotopes are different forms of the same element. They have:
- The same number of protons (same atomic number)
- Different numbers of neutrons (different mass numbers)
- Identical chemical properties (chemistry depends on electrons, not neutrons)
- Different physical properties (mass affects density, diffusion rates, etc.)
Examples of isotopes:
- Carbon-12 (¹²₆C) has 6 protons, 6 neutrons, 6 electrons
- Carbon-13 (¹³₆C) has 6 protons, 7 neutrons, 6 electrons
- Carbon-14 (¹⁴₆C) has 6 protons, 8 neutrons, 6 electrons
Some isotopes are stable, whilst others are radioactive (unstable). Radioactive isotopes have unstable nuclei that decay spontaneously, releasing radiation.
Development of the atomic model
Scientific understanding of atomic structure has evolved significantly over time as new evidence became available.
Dalton's model (early 1800s)
John Dalton proposed that atoms were tiny, solid spheres that could not be divided. Different elements were made of different types of atoms. This model explained many observations but was later proven incomplete.
Thomson's plum pudding model (1897)
J.J. Thomson discovered electrons. He proposed that atoms were spheres of positive charge with tiny negative electrons embedded in them, like plums in a pudding. This model suggested the atom could be divided into smaller parts.
Rutherford's nuclear model (1909-1911)
Ernest Rutherford, along with Hans Geiger and Ernest Marsden, conducted the famous alpha particle scattering experiment (also called the gold foil experiment).
The experiment involved:
- Firing alpha particles (positively charged) at a very thin sheet of gold foil
- Detecting where the alpha particles went using a fluorescent screen
Expected results (based on plum pudding model):
All alpha particles should pass straight through or be deflected only slightly, as the positive charge was thought to be spread throughout the atom.
Actual results:
- Most alpha particles passed straight through (showing atoms are mostly empty space)
- Some alpha particles were deflected through small angles
- A very small number (about 1 in 8000) were deflected back through angles greater than 90°
Conclusions:
Rutherford concluded that:
- The atom's positive charge and most of its mass are concentrated in a tiny nucleus
- The nucleus is extremely small compared to the atom
- Most of the atom is empty space
- Electrons orbit the nucleus at a distance
This experiment provided the evidence that led to the nuclear model of the atom.
Bohr's model (1913)
Niels Bohr refined Rutherford's model by suggesting that electrons orbit the nucleus at specific distances in fixed energy levels or shells. Electrons cannot exist between these energy levels. This explained why atoms only emit or absorb light at specific frequencies.
Later developments
Further experiments revealed that the nucleus contains protons (discovered by Rutherford in 1919) and neutrons (discovered by James Chadwick in 1932). The modern model represents electrons as existing in orbitals (regions of space where they are likely to be found), but the GCSE-level model treats electrons as existing in shells around the nucleus.
Electron arrangement
Electrons occupy energy levels (shells) around the nucleus. The shells closest to the nucleus have lower energy. Electrons fill the lowest available energy levels first.
The rules for GCSE are:
- First shell holds maximum 2 electrons
- Second shell holds maximum 8 electrons
- Third shell holds maximum 8 electrons (at GCSE level)
Examples:
- Carbon (6 electrons): 2, 4
- Oxygen (8 electrons): 2, 6
- Sodium (11 electrons): 2, 8, 1
- Chlorine (17 electrons): 2, 8, 7
When atoms absorb electromagnetic radiation, electrons move to higher energy levels (further from the nucleus). When electrons move to lower energy levels, electromagnetic radiation is emitted. The energy of the radiation equals the energy difference between the levels.
Size and scale
Understanding the relative sizes in atomic structure is essential:
Radius of an atom: approximately 1 × 10⁻¹⁰ m (0.1 nm)
Radius of a nucleus: approximately 1 × 10⁻¹⁴ m
Ratio: The nucleus is about 1/10,000 (or 10⁻⁴) the size of the atom
To put this in perspective: if an atom were the size of a football stadium (100 m across), the nucleus would be the size of a marble (1 cm) in the centre.
Despite occupying such a tiny volume, the nucleus contains over 99.9% of the atom's mass because:
- Protons and neutrons have relative mass of 1
- Electrons have relative mass of approximately 1/2000
- There are roughly equal numbers of protons and neutrons in most stable nuclei
Worked examples
Example 1: Calculating subatomic particles
Question: A chlorine atom is represented as ³⁵₁₇Cl. Calculate the number of protons, neutrons and electrons in this atom. [3 marks]
Answer:
Number of protons = atomic number = 17 ✓
Number of neutrons = mass number - atomic number = 35 - 17 = 18 ✓
Number of electrons = number of protons = 17 (atom is neutral) ✓
Mark scheme: 1 mark for each correct answer with working.
Example 2: Isotopes
Question: Uranium has two main isotopes: uranium-235 and uranium-238. Both isotopes have 92 protons.
(a) Define the term isotope. [2 marks]
(b) Calculate how many neutrons are in uranium-238. [2 marks]
(c) Explain why these isotopes have identical chemical properties. [2 marks]
Answer:
(a) Isotopes are atoms of the same element ✓ with the same number of protons but different numbers of neutrons (or different mass numbers). ✓
(b) Neutrons = mass number - atomic number ✓ = 238 - 92 = 146 neutrons ✓
(c) Chemical properties depend on the number of electrons ✓. Both isotopes have the same number of electrons (92) because they have the same number of protons, so they react in the same way. ✓
Mark scheme: Clear definitions needed for part (a). Part (b) requires working. Part (c) needs explanation linking electrons to chemical behaviour.
Example 3: Alpha scattering experiment
Question: Describe the observations made in Rutherford's alpha particle scattering experiment and explain what these observations showed about the structure of the atom. [6 marks]
Answer:
Observations:
Most alpha particles passed straight through the gold foil ✓
Some alpha particles were deflected through small angles ✓
A very small number were deflected through angles greater than 90° (or bounced back) ✓
Conclusions:
Most particles passing through shows that atoms are mostly empty space ✓
Some deflections show that there is a positive charge in the atom (which repels the positive alpha particles) ✓
Large deflections show that the positive charge and mass are concentrated in a very small region called the nucleus ✓
Mark scheme: 3 marks for observations, 3 marks for explanations. Links between observations and conclusions must be clear.
Common mistakes and how to avoid them
Confusing mass number and atomic number: Remember that atomic number (bottom) is the number of protons, whilst mass number (top) is protons plus neutrons. The atomic number defines the element.
Forgetting electrons have negligible mass: Students often add electrons when calculating atomic mass. The relative mass of an electron is 1/2000, which rounds to 0 at GCSE level.
Thinking isotopes have different numbers of electrons: Isotopes of the same element have different numbers of neutrons, not electrons. In neutral atoms, electron number always equals proton number.
Describing the plum pudding model incorrectly: The positive charge was the "pudding" (spread throughout) with negative electrons embedded in it, not the other way around.
Misinterpreting the alpha scattering results: The key observation is that only a small number of alpha particles were deflected through large angles. This shows the nucleus is small. If it were large, many more would be deflected.
Using incorrect units for atomic radius: Always express atomic radius in metres using standard form (1 × 10⁻¹⁰ m), not just "very small" or informal units.
Exam technique for "Atomic structure: the atom and the nucleus"
"Describe" questions: Provide observations or features without explanation. In the context of the alpha scattering experiment, state what was observed (angles of deflection) without explaining why.
"Explain" questions: Link observations to underlying scientific reasons. For atomic structure, connect observations to properties like charge, mass, or distribution. Aim for 1-2 marks per clear point made.
Calculation questions: Always show your working, even for simple subtractions. Write the formula (neutrons = mass number - atomic number) before substituting values. Include units where appropriate.
Quality of Written Communication (QWC) questions: Some 6-mark questions require extended responses. Use scientific terminology correctly, structure your answer logically (observations, then conclusions), and check spelling of key terms like "nucleus" and "electron".
Quick revision summary
Atoms consist of a tiny, dense nucleus containing protons and neutrons, surrounded by electrons in shells. The nucleus has a radius of about 1 × 10⁻¹⁴ m, whilst the atom has a radius of about 1 × 10⁻¹⁰ m. Atomic number equals the number of protons; mass number equals protons plus neutrons. Isotopes have the same atomic number but different mass numbers. Rutherford's alpha scattering experiment provided evidence for the nuclear model by showing most alpha particles passed through gold foil, but some were deflected, proving the existence of a small, dense, positive nucleus.